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AS CHEMISTRY Unit 2 RevisionRevision BookletBooklet
2.1 Shapes of molecules and ions; 2.2 Intermediate bonding – Intermolecular forces; 2.3 REDOX; 2.4 Periodic Table - Groups 1, 2 and 7; 2.5 Kinetics; 2.6 Equilibria; 2.7 Organic Chemistry 2 - Alcohols, Haloalkanes; 2.8 Mass spec and IR spectroscopy; 2.9 Green Chemistry. 2 3
2.1 – Shapes of Molecules and Ions
Draw the shapes of the following, giving the bond angles and the name of the shapes.
BeCl2, BCl3,
CH4, HCl,
+ NH3, NH4 ,
H2O, CO2,
PCl5, SF6
Explain the values for the bond angles in BF3, CH4, NH3 and H2O . 4
How do chemists find bond energies and bond lengths?
Draw a diagram to show the structure of the two carbon allotropes, diamond and graphite.
What are carbon nanotubes?
Give a use for carbon nanotubes. 5
2.2 - Intermediate bonding and bond polarity
What is meant by the term electronegativity?
What happens to the electrons in a covalent bond if two atoms are not the same?
Show the dipole moments in the following molecules? O O O H F H H C C H 3 C C H 3 F F
H H H F F H C C
H O H
Explain why CCl4 isn’t a polar molecule.
Describe an experiment that would show whether a molecule is polar or not.
Draw a diagram to show the electron density map of a polar covalent molecule.
Draw an electron density map showing the polarisation in MgCl2 . 6
Intermolecular forces
What are the three common examples of intermolecular force and how do they arise?
1)
2)
3)
Explain why iodine is soluble in hexane but not soluble in water.
Explain why glucose is soluble in water but not soluble in hexane.
Explain why hexan-1-ol has a higher boiling point than ethanol.
Explain why ethanol has a higher boiling point than ethane. 7
Describe three unusual properties of water caused by hydrogen bonding.
How are Lattice enthalpies, solution enthalpies and hydration enthalpies related?
Draw a diagram to show the hydration of a sodium ion.
Explain why hydration enthalpies are always negative. 8
2.3 - REDOX
What is a redox reaction?
What is the oxidation number of sulphur in the following substances? Na2S SO2 SO42-
What has been oxidized and reduced in the following redox reaction?
VO+ + 2H+ VO3+ + H2
Write half equations for each of the following conversions.
- a) Br Br2
2- b) SO4 SO2
2- 3+ c) Cr2O7 Cr
Use the following half equations to create the balanced equations for the reactions below. - + - 2+ MnO4 + 8 H + 5 e Mn + 4 H2O Fe3+ + e- Fe2+ + - 2 H + O2 + 2 e H2O2
2+ + - a) oxidation of Fe by H /MnO4
+ - b) oxidation of H2O2 by H /MnO4 9
2.4 - Periodic Table - Group 2
Explain the following trends Mg Ca Sr Ba 1st IE kJ/mol 738 590 550 530
Mg Ca Sr Ba Mpt oC 649 839 769 725
Write a balanced equation for the reaction of calcium with the following; Oxygen
Water
Steam
Chlorine
Write equations for the following reactions;
The addition of water to calcium oxide.
The test for carbon dioxide.
The reaction between calcium hydroxide and hydrochloric acid. 10
Thermal stability and solubility
Write an equation for the thermal decomposition of calcium carbonate.
Explain in terms of lattice enthalpies why the carbonates of Group 2 metals become more stable to heat down the group.
Write an equation for the thermal decomposition of the Group 2 nitrate, calcium nitrate.
Write an equation for the thermal decomposition of the Group 1 nitrate, sodium nitrate.
Explain why Group 1 and 2 nitrates decompose in different ways.
Compound solubility Describe the solubility trends of the following Group 2 compounds.
Sulphates
Hydroxides
Carbonates
Nitrates
Why doesn’t barium oxide react fully with sulphuric acid? 11
Explain why the solubility of the Group 2 sulphates decreases down the group.
Explain why the solubility of the Group 2 hydroxides increases down the group.
Complete the table below showing the flame colours of Group 1 and 2 ions.
Metal Ion Flame Colour Lithium
Sodium
Potassium
Calcium
Strontium
Barium
What causes the metal ion to have a flame colour?
Why can’t a flame test be used on a mixture? 12
2.4 - Periodic Table - Group 7
Complete the following table;
Element Mpt oC Bpt oC State at Colour of Colour of rtp element vapour Fluorine -220 -188 Chlorine -101 -34 Bromine -7 59 Iodine 114 184 Astatine
What is the order of oxidant strength of the halogens?
Explain this order in terms of electronic structure.
Write a balanced equation for the reaction between chlorine and Fe2+ ions.
What are the common properties of the hydrogen halides?
What are the products when sodium chloride is reacted with concentrated sulphuric acid?
What are the products when sodium bromide is reacted with concentrated sulphuric acid?
What are the products when sodium iodide is reacted with concentrated sulphuric acid? 13
Complete the oxidation number chart below;
Formula Oxidation no of halogen Name of compound KClO4 +7 ...... KClO3 +5 ...... KClO2 ...... KClO +1 ...... Cl2 ...... KCl -1 ......
What happens to the chlorine in the reaction below?
Cl2 + 2NaOH NaCl + NaClO + H2O
What is the common, household, name for NaClO?
If the alkali had been hot the following reaction occurs;
Cl2 + NaOH NaCl + NaClO3 + H2O
Balance this redox equation.
Complete the following table; Action of Halide + Silver nitrate + Ammonia (dil) + Ammonia (conc.) Sunlight
KCl ......
KBr ......
KI ...... 14
Describe the titration used to find the concentration of Iodine in a solution.
A 0.800g of a contaminated sample of potassium iodate, KIO3, was dissolved in 250cm3 of solution. 25cm3 of this solution was added to an excess of potassium iodide and dilute sulphuric acid.
5KI + KIO3 + 3H2SO4 3I2 + 3H2O + 3K2SO4
The mixture required 21.60cm3 of 0.1 mol dm-3 sodium thiosulphate solution to remove the iodine released.
2- - 2- 2S2O3 + I2 2I + S4O6
Calculate the percentage purity of the potassium iodate. 15
2.5 - Kinetics
What factors affect the rate of reaction?
Explain the effect of these factors, quantitatively.
Use a Maxwell-Boltzmann distribution to explain why the rate of a reaction increases with temperature?
What is meant by the activation energy for a reaction?
What is a catalyst?
What are Homo and Heterogeneous Catalysis? (give an example of each one). 16
2.6 - Equilibria
What is meant by the term equilibrium?
What is Le Chateliers Principle?
How can the conditions in the Haber process be controlled to obtain an optimum yield of NH3?
3H2 + N2 2NH3 H = -92.4 kJ/mol
What are the conditions used in the Haber Process?
Temp ...... Pressure ...... Catalyst ......
Explain why these conditions are chosen.
Explain what would be observed if the following changes were made to the equilibrium below. Cl2(g) + ICl(l) ICl3(s) Pale green Brown yellow
More chlorine is added.
The pressure is decreased.
A catalyst is added. 17
2.7 – Organic Chemistry – Halogenoalkanes
Define what is meant be a primary and secondary halogenoalkane.
Describe an experiment to determine if a chloroalkane is more or less reactive than an iodoalkane?
Which is more reactive, a chloroalkane or an iodoalkane? Explain your answer.
Describe how you would make 2-chlorobutane from butan-2-ol.
Give two uses of halogenoalkanes. 18
Draw and name the products formed from the following reactions; 1) 2-Chloropropane and aqueous sodium hydroxide.
2) 2-Chlorobutane and ethanolic potassium hydroxide.
3) 1-Bromobutane and alcoholic ammonia.
Do any of the reactions above produce isomeric products? Explain you answer.
Give the mechanism for the reaction between 2-chloropropane and aqueous sodium hydroxide.
What name is given to the mechanism above. 19
2.7 – Organic Chemistry – Alcohols Define what is meant be a primary and secondary alcohol.
Explain why short chained alcohols are soluble in water.
Draw and name the products formed from the following reactions; 1) Propan-1-ol and acidified potassium dichromate (Give the colour change)
2) Hexan-2-ol and ethanolic sodium hydroxide.
3) Hexan-3-ol and phosphorous pentachloride.
4) Propan-2-ol and acidified potassium dichromate (Give the colour change).
Do any of the reactions above produce isomeric products? Explain. 20
Describe how you would make propanal from propan-1-ol.
Describe how you modify the experiment to make propanoic acid from propan-1-ol.
Why are organic reactions heated under reflux?
Describe how IR spectroscopy could be used to prove that propanoic acid had been made.
Draw the skeletal formulae of the following molecules.
Propanal Propanoic acid Propanone
State the type of homologous series that each of the above molecules belong to.
Describe the chemical test for an O-H functional group. 21
2.8 – Mass spec and IR spectroscopy
Describe how the following isomers can be identified using Mass spec and IR spectroscopy.
O H 2 H2 C C O C C OH C CH3 H2
O
Which type of molecules can absorb infrared radiation?
Why are H2O, CO2, CH4 and NO molecules greenhouse gases, whilst O2 and N2 are not? 22
2.9 – Green Chemistry
State five ways in which the chemical industry is becoming more sustainable ('greener').
Explain what is meant by the following terms;
‘natural climate change’
‘anthropogenic climate change’
'carbon neutrality'
'carbon footprint'
Give an example of a carbon neutral fuel. 23
Explain the science community's reasons for recommending that CFC’s are no longer used.
Describe how ozone, O3, absorbs UV radiation.
Describe how ozone, O3, depletion occurs. Including the mechanism for the reaction. 24